Final answer:
The first ionization energy of potassium is lower than argon's because potassium’s valence electron is further from the nucleus in a higher energy level and experiences more electron shielding. This reduces the effective nuclear charge and makes it easier to remove that electron, despite potassium's higher nuclear charge.
Step-by-step explanation:
The first ionization energy of an element is the energy required to remove the most loosely held electron from an atom in its gaseous state. In the case of potassium (K), the first ionization energy is lower than that of argon (Ar), despite potassium having a higher nuclear charge. The reason can be explained by considering which quantum shell the outer electron occupies and the electron shielding effect. The valence electron in potassium (K) is in the 4s orbital, which is in the fourth energy level. Because higher energy levels are farther away from the nucleus, the electrostatic attraction to the nucleus is weaker. This effect is compounded by electron shielding, where the inner electrons block the valence electron from the full effect of the nuclear charge, making it easier to remove. On the other hand, argon (Ar) has its outermost electron in the 3p orbital, but as a noble gas, its electron configuration is especially stable. Additionally, the electron shielding in argon is less pronounced than in potassium because argon’s electrons are in a lower energy level (closer to the nucleus) where there is less shielding by inner electrons. This makes the first ionization energy of argon much higher.
Therefore, despite the higher nuclear charge in potassium, the combined effects of the outer electron being in a higher quantum shell and experiencing more electron shielding results in a much lower first ionization energy compared to argon.